Bond Polarities and Molecular Polarity

Last updated: 16 August 2026

Diagnose → Learn → Master: Every study-note section includes 15 free MCQs — start with 5 Diagnostic MCQs to identify gaps in your understanding, then complete 10 Mastery MCQs near the end to check what you have learned.

When two atoms form a covalent bond, we often say that they share a pair of electrons. But “sharing” does not necessarily mean that both atoms pull equally strongly on that electron pair.

If one atom attracts the bonding electrons more strongly, electron density shifts towards that atom. The bond then develops a slightly negative end and a slightly positive end. This is what we mean by bond polarity.

The next question is more subtle: does that make the whole molecule polar? Not always. To decide this, we must combine bond polarity with the three-dimensional shape of the molecule and determine whether the individual bond dipoles cancel.

Diagnostic Quiz

What Do You Already Know About Polarity?

Try these five questions before reading the chapter. They check whether you can connect electronegativity, bond polarity and molecular shape.

5 quick questions
Focus on the key ideas behind polarity
Every answer option is explained

Which statement best describes electronegativity?

The energy required to remove an electron from an isolated gaseous atom

This describes ionisation energy, not electronegativity. Electronegativity concerns electrons that are already being shared in a covalent bond.

The ability of an atom in a covalent bond to attract the shared pair of electrons towards itself

Correct. Electronegativity compares how strongly bonded atoms attract the shared electron pair.

The number of electrons in the outermost occupied shell

The number of outer-shell electrons affects chemical behaviour, but it is not the definition of electronegativity.

The tendency of an atom to lose all of its valence electrons

Electronegativity describes attraction for a shared electron pair, not simply a tendency to lose electrons.

In an H—Cl bond, chlorine is more electronegative than hydrogen. Which description is correct?

H is δ− and Cl is δ+

This reverses the partial charges. The bonding electrons are attracted more strongly towards chlorine.

Both H and Cl carry full ionic charges

H—Cl contains a polar covalent bond. The electrons are shared unequally, but there has not been complete electron transfer to form H⁺ and Cl⁻.

H and Cl have no partial charges because the bond is covalent

Covalent bonding does not require equal sharing. A difference in electronegativity can produce partial charges within a covalent bond.

H is δ+ and Cl is δ−

Correct. Chlorine attracts the shared electron pair more strongly, so electron density shifts towards chlorine.

Carbon dioxide contains two polar C=O bonds. Why is a CO₂ molecule nevertheless non-polar?

CO₂ is linear, so its two equal bond dipoles act in opposite directions and cancel

Correct. Molecular polarity depends on the vector sum of the bond dipoles. In linear CO₂, the two equal C=O dipoles cancel completely.

The C=O bonds become non-polar when two of them are present

Each C=O bond remains polar. It is the overall molecular dipole that becomes zero because the two bond dipoles cancel.

Carbon and oxygen have identical electronegativities

Oxygen is more electronegative than carbon, so each C=O bond is polar.

CO₂ contains no shared electron pairs

CO₂ contains covalent C=O bonds formed by shared electrons. The key issue is cancellation of the bond dipoles.

Why is H₂O polar whereas CO₂ is non-polar?

Only H₂O contains covalent bonds

Both H₂O and CO₂ are covalent molecules. The difference in overall polarity comes from molecular shape and dipole cancellation.

Oxygen is electronegative in H₂O but not in CO₂

Oxygen is strongly electronegative in both molecules. Both O—H and C=O bonds are polar.

H₂O is bent, so its O—H bond dipoles do not cancel completely

Correct. The O—H dipoles act at an angle and give a non-zero resultant. CO₂ is linear, so its equal C=O dipoles cancel.

H₂O has an overall ionic charge

A water molecule is electrically neutral overall. Being polar means it has an uneven distribution of charge, not a net ionic charge.

Which statement about molecular polarity is correct?

Every molecule containing a polar bond must be polar

Polar bonds can cancel one another. CO₂, BF₃ and CCl₄ are important examples of molecules that contain polar bonds but are non-polar overall.

Both bond polarity and molecular shape must be considered

Correct. Bond polarity tells you the individual bond dipoles, while molecular shape determines how those dipoles combine.

Molecular polarity depends only on the number of atoms present

The number of atoms alone is not enough. Their electronegativities and three-dimensional arrangement are important.

A symmetrical molecule must contain only non-polar bonds

A symmetrical molecule may contain polar bonds whose dipoles cancel. BF₃ and CCl₄ are examples.

Electronegativity: where bond polarity begins

To understand why some covalent bonds are polar, begin with electronegativity.

Electronegativity is the ability of an atom in a covalent bond to attract the shared pair of electrons towards itself.

Electronegativity is about attraction for a shared pair of electrons.

We are comparing how strongly the two atoms at the ends of a covalent bond pull on the same bonding electron pair.

What happens when the atoms have the same electronegativity?

Consider a bond between two identical atoms, such as H—H or Cl—Cl.

Since the two atoms are identical, they have the same electronegativity. Neither atom attracts the bonding electrons more strongly than the other, so the electron density is shared equally.

The bond therefore has no permanent separation of charge and is described as a non-polar covalent bond.

What happens when the atoms have different electronegativities?

Now consider H—Cl. Chlorine is more electronegative than hydrogen, so chlorine attracts the shared pair of electrons more strongly.

Electron density is therefore greater near chlorine and lower near hydrogen. The bond develops two partial charges:

H^δ+—Cl^δ−

The chlorine atom carries a partial negative charge, δ−, while the hydrogen atom carries a partial positive charge, δ+.

The bond now has two poles and is described as a polar covalent bond.

General electronegativity trends

You do not normally need to memorise a long table of electronegativity values, but the general Periodic Table pattern is useful.

Movement in the Periodic Table	General electronegativity trend
Across a period from left to right	Electronegativity generally increases
Down a group	Electronegativity generally decreases

Fluorine is the most electronegative element. Oxygen and nitrogen are also highly electronegative.

The important idea for bond polarity is the difference in electronegativity between the two bonded atoms.

The greater the electronegativity difference, the more polar the bond tends to be.

A larger difference means that the shared electrons are displaced more strongly towards one end of the bond, producing a greater separation of partial charge.

Bond polarity and bond dipoles

A polar covalent bond contains an uneven distribution of electron density. The resulting separation of partial positive and partial negative charge is called a bond dipole.

For H—Cl:

H^δ+—Cl^δ−

Chlorine is the negative end of the dipole because it attracts the bonding electrons more strongly.

δ+ and δ− are partial charges

The symbol δ is important. It tells us that these are only partial charges.

We are not saying that an electron has been transferred completely from hydrogen to chlorine. HCl remains a covalent molecule in which the bonding electron pair is shared, although it is shared unequally.

H^δ+—Cl^δ−: the charges are partial because the electrons are still being shared in a covalent bond.

Na⁺ and Cl⁻: these are full ionic charges resulting from electron transfer.

Representing the direction of a bond dipole

When a dipole arrow is used, the arrowhead points towards the more electronegative, δ− end of the bond.

Therefore, for H—Cl, the dipole is directed towards chlorine.

Follow the electrons.

Ask which atom attracts the bonding pair more strongly. That atom is the δ− end, and the bond dipole points towards it.

From bond polarity to molecular polarity

This is the point where bond polarity and molecular polarity must be separated carefully.

A molecule may contain several polar bonds. Each bond has its own dipole, with its own magnitude and direction.

To determine whether the whole molecule is polar, we must find the overall result of all these individual bond dipoles.

Molecular dipole moment = vector sum of the individual bond dipoles

Bond dipoles are therefore treated as vectors. Their directions matter as well as their sizes.

When do bond dipoles cancel?

If equal bond dipoles are arranged symmetrically, they may cancel completely. The molecule then has no net dipole moment and is non-polar.

If the dipoles do not cancel completely, there is a resultant dipole. The molecule then has a net dipole moment and is polar.

Polar bonds do not automatically give a polar molecule.

First identify the bond dipoles. Then determine the molecular shape and ask whether those dipoles cancel as vectors.

Bond polarity: asks whether one particular covalent bond has unequal sharing of electron density.

Molecular polarity: asks whether all the dipoles in the complete molecule combine to give a non-zero resultant.

Carbon dioxide: polar bonds but a non-polar molecule

Carbon dioxide, CO₂, is one of the most important examples because it shows why bond polarity alone is not enough.

Each C=O bond is polar because oxygen is more electronegative than carbon. Electron density is pulled towards each oxygen atom.

However, CO₂ is linear.

O=C=O

The two C=O bond dipoles have equal magnitudes and point in exactly opposite directions.

Their vector sum is therefore zero.

Conclusion: CO₂ contains polar C=O bonds, but the molecule is non-polar because the two equal bond dipoles cancel completely in its linear shape.

Water: polar bonds and a polar molecule

Water, H₂O, provides a useful contrast with carbon dioxide.

Each O—H bond is polar because oxygen is more electronegative than hydrogen. The bond dipoles point towards oxygen.

But H₂O is not linear. It is bent.

The two O—H bond dipoles therefore act at an angle rather than in exactly opposite directions.

They cannot cancel completely. Their vector sum gives a resultant dipole directed towards the oxygen side of the molecule.

Conclusion: H₂O is polar because its bent shape prevents the O—H bond dipoles from cancelling.

CO₂: polar bonds + linear shape + complete cancellation = non-polar molecule.

H₂O: polar bonds + bent shape + incomplete cancellation = polar molecule.

Boron trifluoride: cancellation in a trigonal planar molecule

BF₃ contains three B—F bonds. Fluorine is much more electronegative than boron, so each B—F bond is polar.

BF₃ is trigonal planar. The three B—F bonds are arranged evenly around the central boron atom, with bond angles of 120°.

The three identical bond dipoles therefore point symmetrically away from the central atom.

When the three vectors are added, they cancel completely.

Conclusion: BF₃ contains three polar B—F bonds but is non-polar because the three equal bond dipoles cancel in the trigonal planar arrangement.

Tetrachloromethane: cancellation in three dimensions

Tetrachloromethane, CCl₄, is another important example.

Each C—Cl bond is polar because chlorine is more electronegative than carbon. Yet CCl₄ is non-polar overall.

The carbon atom is surrounded by four chlorine atoms in a tetrahedral arrangement.

The four identical C—Cl bond dipoles point symmetrically in three-dimensional space and cancel completely.

Conclusion: CCl₄ is non-polar because its four identical C—Cl bond dipoles cancel in the symmetrical tetrahedral arrangement.

Think in three dimensions.

A displayed formula drawn on a flat page may not make dipole cancellation obvious. Molecular polarity must be judged using the actual three-dimensional molecular shape.

What happens if one chlorine is replaced by hydrogen?

Now compare CCl₄ with CHCl₃.

The arrangement around carbon remains tetrahedral, but the four surrounding atoms are no longer identical. One position is occupied by hydrogen rather than chlorine.

The bond dipoles therefore no longer form four equal vectors arranged symmetrically. They do not cancel completely.

Conclusion: CHCl₃ is polar because replacing one Cl atom with H destroys the symmetry required for complete dipole cancellation.

Do lone pairs automatically make a molecule polar?

Lone pairs often lead to polar molecules because they can produce an asymmetrical molecular shape.

For example, H₂O is bent because oxygen has two bonding pairs and two lone pairs. NH₃ is trigonal pyramidal because nitrogen has three bonding pairs and one lone pair.

In both cases, the bond dipoles do not cancel completely, so the molecules are polar.

However, the statement “a molecule is polar because the central atom has a lone pair” is not a sufficiently general rule.

Do not use the presence of a lone pair as the final test for polarity.

Lone pairs matter because they can affect molecular shape. The final decision must still be based on the overall arrangement of the bond dipoles and whether their vector sum is zero.

A reliable method for determining molecular polarity

When an examination question asks whether a molecule is polar, use the same reasoning sequence every time.

Step 1: identify the polar bonds

Compare the electronegativities of the bonded atoms.

For each polar bond, identify:

the less electronegative atom as δ+;
the more electronegative atom as δ−; and
the direction of the bond dipole towards the more electronegative atom.

Step 2: determine the molecular shape

Use the electron-pair arrangement around the central atom to determine the actual molecular geometry.

Typical examples in this topic include:

linear CO₂;
bent H₂O;
trigonal planar BF₃;
trigonal pyramidal NH₃; and
tetrahedral CCl₄ and CHCl₃.

Step 3: combine the bond dipoles as vectors

Ask whether the dipoles have equal magnitudes and are arranged so that they cancel.

Complete cancellation → no net dipole moment → non-polar molecule

Incomplete cancellation → net dipole moment → polar molecule

Step 4: write the complete explanation

A strong explanation should not simply say that a molecule is “symmetrical” or “asymmetrical”.

Explain what happens to the bond dipoles.

Example: CO₂ is linear. The two C=O bonds are polar, but their equal bond dipoles act in opposite directions and cancel completely. CO₂ therefore has no net dipole moment and is non-polar.

Comparing common molecular shapes and polarity

Molecule	Shape	Bond polarity	Do the dipoles cancel?	Molecular polarity
CO₂	Linear	C=O bonds are polar	Yes	Non-polar
H₂O	Bent	O—H bonds are polar	No	Polar
BF₃	Trigonal planar	B—F bonds are polar	Yes	Non-polar
NH₃	Trigonal pyramidal	N—H bonds are polar	No	Polar
CCl₄	Tetrahedral	C—Cl bonds are polar	Yes	Non-polar
CHCl₃	Tetrahedral	Bonds have different dipole contributions	No	Polar

Symmetry is useful only when the bond dipoles are equivalent. BF₃ and CCl₄ have identical surrounding atoms arranged symmetrically, so their bond dipoles cancel.

Shape alone is not enough. CCl₄ and CHCl₃ are both tetrahedral, but CHCl₃ has different surrounding atoms, so its bond dipoles do not cancel completely.

Common mistakes with bond and molecular polarity

Common mistake	Why it is wrong	Correct approach
Assuming that every covalent bond shares electrons equally	Atoms with different electronegativities attract the shared pair by different amounts	Use electronegativity to decide whether the bond is polar
Putting δ− on the less electronegative atom	Electron density is displaced towards the more electronegative atom	Place δ− on the more electronegative atom and δ+ on the other atom
Treating δ+ and δ− as full ionic charges	The electrons are still shared in a polar covalent bond	Describe δ+ and δ− as partial charges
Assuming that a molecule containing polar bonds must be polar	Bond dipoles may cancel because of molecular symmetry	Determine the shape and add the dipoles as vectors
Saying that CO₂ is non-polar because C=O bonds are non-polar	C=O bonds are polar	State that the equal C=O dipoles cancel because CO₂ is linear
Saying that BF₃ is polar because B—F bonds are polar	The three equal B—F dipoles cancel in the trigonal planar arrangement	Consider the vector sum of the three bond dipoles
Assuming that every tetrahedral molecule is non-polar	Cancellation requires an appropriate symmetrical arrangement of equivalent bond dipoles	Compare non-polar CCl₄ with polar CHCl₃
Judging polarity from a flat displayed formula	The actual molecule is three-dimensional	Use the VSEPR molecular shape before analysing dipole cancellation
Saying that a lone pair automatically makes a molecule polar	Polarity depends on the overall vector sum of the molecular dipoles	Use lone pairs to determine shape, then test whether the dipoles cancel

Bond-polarity and molecular-polarity checklist

Check the bonds

Compare the electronegativities of the bonded atoms.
Identify which bonds are polar.
Put δ− on the more electronegative atom.
Put δ+ on the less electronegative atom.

Find the shape

Use VSEPR theory to determine the molecular shape.
Think about the actual three-dimensional arrangement.
Check whether the surrounding atoms are identical.
Do not judge from the flat displayed formula alone.

Combine the dipoles

Treat each bond dipole as a vector.
Ask whether equal dipoles cancel completely.
Zero resultant means a non-polar molecule.
A non-zero resultant means a polar molecule.

Explain your conclusion

State which bonds are polar.
Name the molecular shape.
State whether the dipoles cancel.
Conclude whether the molecule is polar or non-polar.

Mastery Quiz

Bond Polarities and Molecular Polarity Mastery Quiz

Now apply the chapter to molecular shapes and unfamiliar comparisons. These questions focus on reasoning from electronegativity, geometry and dipole cancellation rather than simple definition recall.

10 application-focused questions
Shape and dipole reasoning
Every answer option is explained

CO₂ and H₂O both contain polar bonds to oxygen. Which factor best explains why CO₂ is non-polar whereas H₂O is polar?

CO₂ contains double bonds whereas H₂O contains single bonds

Bond order is not the key reason for the difference in molecular polarity. Both molecules contain polar bonds; their different shapes determine whether those dipoles cancel.

Carbon is a non-metal whereas hydrogen is a metal

Hydrogen is also a non-metal. This statement does not explain the difference in dipole cancellation.

CO₂ is linear whereas H₂O is bent

Correct. The equal C=O dipoles in linear CO₂ cancel. The O—H dipoles in bent H₂O act at an angle and give a non-zero resultant.

Oxygen is more electronegative in H₂O than in CO₂

The relevant difference here is molecular geometry, not a change in the identity of oxygen. Oxygen attracts bonding electrons strongly in both.

Each B—F bond in BF₃ is strongly polar. Why does BF₃ have no net dipole moment?

The three equal bond dipoles are arranged symmetrically at 120° and cancel

Correct. BF₃ is trigonal planar. The vector sum of its three identical B—F bond dipoles is zero.

Fluorine becomes less electronegative when three fluorine atoms are present

The B—F bonds remain polar. The absence of a molecular dipole results from cancellation of those dipoles, not from the bonds becoming non-polar.

Boron gains three electrons from fluorine

BF₃ is being considered as a covalent molecule. Its bond polarity arises from unequal sharing, not complete electron transfer to boron.

Any molecule with three bonds must be non-polar

The number of bonds alone does not determine polarity. Their magnitudes and three-dimensional directions matter.

CCl₄ is non-polar, but CHCl₃ is polar. What is the best explanation?

CHCl₃ is not tetrahedral

Carbon has four bonding pairs in CHCl₃, so the arrangement around carbon is tetrahedral.

C—Cl bonds are non-polar in CCl₄ but polar in CHCl₃

C—Cl bonds are polar in both molecules because chlorine is more electronegative than carbon.

CCl₄ has no electrons on chlorine

Chlorine atoms contain electrons and participate in polar C—Cl bonds. Electron absence is not the reason for CCl₄ being non-polar.

The four identical C—Cl dipoles in CCl₄ cancel, whereas replacing one Cl with H removes that symmetry

Correct. CCl₄ has four equivalent bond dipoles in a symmetrical tetrahedral arrangement. CHCl₃ does not, so the resultant dipole is non-zero.

A molecule AX₂ is linear. Both A—X bonds are identical and polar towards X. What can be concluded about the molecule?

It must be polar because it contains two polar bonds

The presence of polar bonds alone is insufficient. Their directions must also be considered.

It is non-polar because the two equal bond dipoles point in opposite directions and cancel

Correct. In a linear AX₂ molecule with identical terminal atoms, equal and opposite bond-dipole vectors give a zero resultant.

The A—X bonds must become non-polar

The individual bonds remain polar. It is the molecular dipole that becomes zero because of vector cancellation.

The molecule must have an overall 2− charge

Partial bond polarity does not imply an overall ionic charge. A neutral molecule can contain polar bonds.

NH₃ has three polar N—H bonds. Which explanation correctly accounts for its molecular polarity?

NH₃ is trigonal planar, so the three dipoles cancel

NH₃ is not trigonal planar. The lone pair on nitrogen gives a trigonal pyramidal molecular shape.

The N—H bonds are non-polar because both atoms are non-metals

Covalent bonds between non-metals can still be polar when the bonded atoms have different electronegativities.

NH₃ is trigonal pyramidal, so the N—H bond dipoles do not cancel completely

Correct. The asymmetrical trigonal pyramidal arrangement gives a non-zero vector sum and hence a net molecular dipole.

Nitrogen has a full 3− ionic charge in NH₃

The N—H bonds are covalent and contain partial charge separation. Nitrogen does not carry a full 3− charge in an NH₃ molecule.

Suppose one F atom in trigonal planar BF₃ were replaced by an atom Y so that the B—Y bond dipole had a different magnitude from the two B—F bond dipoles. What would most likely happen?

The molecule would remain non-polar simply because it is trigonal planar

Trigonal planar geometry guarantees cancellation only when the relevant bond dipoles are arranged symmetrically with suitable equal magnitudes.

Every bond in the molecule would become non-polar

Replacing one atom does not automatically remove the electronegativity differences in the other bonds.

The bond dipoles would no longer cancel completely, so the molecule would have a net dipole

Correct. Once the three dipole vectors are no longer equivalent, their vector sum need not be zero even though the arrangement remains trigonal planar.

The central atom would necessarily acquire a full positive ionic charge

Molecular polarity involves partial charge separation and does not require the formation of a full ion.

Fluorine is more electronegative than chlorine. Which statement is the most reasonable comparison of H—F and H—Cl bond polarity?

H—F is expected to have the greater electronegativity difference and stronger bond polarity

Correct. Fluorine attracts the bonding pair more strongly than chlorine, so the H—F bond has a larger electronegativity difference.

H—Cl must be more polar because chlorine contains more electrons

The total number of electrons in an atom does not by itself determine bond polarity. The relevant comparison is electronegativity.

Both bonds must be non-polar because hydrogen is present in each

Hydrogen has a different electronegativity from both halogens, so unequal sharing occurs in both bonds.

H—F and H—Cl must have identical polarity because both contain one hydrogen atom

The identity and electronegativity of the second bonded atom also matter. Fluorine and chlorine do not attract the shared pair equally strongly.

A student writes H⁺—Cl⁻ to show the polarity of an H—Cl molecule. What is the main problem with this notation?

Chlorine should be positive because it is more electronegative

A more electronegative atom attracts electron density and therefore becomes the partially negative end, not the positive end.

Hydrogen and chlorine cannot form a covalent bond

Hydrogen and chlorine do form a polar covalent H—Cl bond.

The charges should both be zero because covalent bonds can never be polar

Covalent bonds can contain an uneven distribution of electron density and therefore partial charges.

H⁺ and Cl⁻ imply full ionic charges; Hδ+—Clδ− should be used for partial charge separation

Correct. In HCl the electron pair is shared unequally rather than transferred completely, so δ+ and δ− are appropriate.

The two O—H bond dipoles in H₂O point towards oxygen. Why do they not simply cancel each other?

One O—H bond is polar and the other is non-polar

The two O—H bonds are equivalent and both are polar towards oxygen.

The molecule is bent, so the two dipoles are not directed exactly opposite to one another

Correct. Because the dipoles act at an angle, their vector sum is not zero and the molecule has a net dipole moment.

Oxygen has no lone pairs

Oxygen in H₂O has two lone pairs. These contribute to the bent molecular shape.

Dipoles can never cancel in a molecule

Dipole cancellation is exactly why molecules such as CO₂, BF₃ and CCl₄ can be non-polar despite containing polar bonds.

A student states: “Any molecule with a lone pair on its central atom must be polar.” What is the best evaluation of this statement?

It is always correct because a lone pair carries a full negative charge

A lone pair is an electron pair, but it does not give the central atom a full ionic charge merely by being present.

It is always correct because lone pairs make every bond polar

Bond polarity is determined mainly by electronegativity differences between the bonded atoms, not simply by the presence of a lone pair.

It is not a reliable general rule; the molecular geometry and overall dipole cancellation must still be considered

Correct. Lone pairs can change molecular shape, as in H₂O and NH₃, but molecular polarity must ultimately be decided from the overall vector sum of the dipoles.

It is wrong because molecules containing lone pairs are always non-polar

H₂O and NH₃ are familiar counterexamples: both have lone pairs and are polar. The correct approach is to analyse the complete molecular geometry.

Frequently asked questions

Questions students often ask about bond and molecular polarity

These questions connect electronegativity, molecular shape and dipole cancellation so that you can apply the same reasoning to unfamiliar molecules.

Can a molecule have polar bonds but still be non-polar?

Yes. CO₂, BF₃ and CCl₄ are important examples. Their individual bonds are polar, but the bond dipoles are arranged symmetrically and cancel completely. The molecules therefore have no net dipole moment.

Does a non-polar molecule have to contain only non-polar bonds?

No. A non-polar molecule may contain polar bonds if their dipoles cancel. CCl₄ contains four polar C—Cl bonds, but its symmetrical tetrahedral arrangement gives a zero resultant dipole.

Why is CO₂ non-polar but H₂O polar?

CO₂ is linear, so its two equal C=O bond dipoles point in opposite directions and cancel. H₂O is bent, so its two O—H bond dipoles act at an angle and do not cancel completely.

Why is BF₃ non-polar even though fluorine is very electronegative?

Each B—F bond is strongly polar, but BF₃ is trigonal planar. The three identical B—F dipoles are separated symmetrically by 120°, so their vector sum is zero.

Why is CCl₄ non-polar but CHCl₃ polar?

CCl₄ has four identical chlorine atoms arranged tetrahedrally around carbon, so its four equivalent C—Cl bond dipoles cancel. In CHCl₃, one chlorine is replaced by hydrogen. The bond dipoles are no longer equivalent and therefore do not cancel completely.

Does a symmetrical shape always make a molecule non-polar?

Not by shape alone. The bond dipoles must also have suitable equal magnitudes. CCl₄ is non-polar because its four identical surrounding atoms give four equivalent dipoles. A tetrahedral molecule with different surrounding atoms may still be polar.

Does a lone pair always make a molecule polar?

No. Lone pairs are important because they can affect molecular geometry, but the final test is still whether the overall dipoles cancel. In familiar examples such as H₂O and NH₃, the lone pairs lead to shapes in which the bond dipoles do not cancel.

What is the difference between δ− and a negative ionic charge?

δ− represents a partial negative charge caused by an unequal distribution of shared electrons. A full negative ionic charge, such as Cl⁻, means that the species has gained enough electron charge to have a net charge of −1.

Why must bond dipoles be treated as vectors?

A bond dipole has both magnitude and direction. Equal dipoles cancel only if their directions are arranged appropriately. This is why molecular shape is essential when determining molecular polarity.

What should I write in an examination when asked whether a molecule is polar?

Identify the polar bonds, state the molecular shape, and explain whether the bond dipoles cancel. Then state whether the molecule has a net dipole moment and conclude whether it is polar or non-polar.

Bond polarity begins with electronegativity: when two bonded atoms attract the shared pair of electrons unequally, partial charges and a bond dipole develop. Molecular polarity then adds one crucial idea — shape. The individual bond dipoles must be combined as vectors. If they cancel completely, as in CO₂, BF₃ and CCl₄, the molecule is non-polar. If they do not cancel, as in H₂O, NH₃ and CHCl₃, the molecule has a net dipole moment and is polar. Once you separate these two stages — bond polarity first, molecular geometry second — polarity becomes a reasoning problem rather than a list of molecules to memorise.